Reaction Kinetics
Reaction kinetics is the study of how fast chemical reactions go and what controls their speed. The PMDC MDCAT 2026 syllabus expects you to define rate, write rate laws, identify the order, describe the role of the rate constant, and use activation energy and the activated complex to explain why temperature and catalysts matter. This chapter typically yields 2-3 MCQs.
Chemical Kinetics
Chemical kinetics is the branch of chemistry that deals with the rate of a reaction and the mechanism by which it proceeds. Rate measures how fast reactants are consumed (or products formed) per unit time:
Rate = −Δ[Reactant] / Δt = +Δ[Product] / Δt
Units of rate: mol dm−3 s−1. For aA + bB → cC + dD, rate = −(1/a) d[A]/dt = −(1/b) d[B]/dt = (1/c) d[C]/dt = (1/d) d[D]/dt.
Rate law (rate equation)
For a reaction aA + bB → products, the experimental rate law has the form:
Rate = k [A]m [B]n
where m and n are determined experimentally (not from the balanced equation), k is the rate constant, and m + n is the overall order.
Rate Constant
The rate constant k (also called specific rate constant) is the proportionality constant in the rate law. It is numerically equal to the rate when all reactant concentrations are 1 mol dm−3.
- k depends on temperature and the catalyst, but is independent of concentration.
- Units of k depend on the overall order:
- Zero order: mol dm−3 s−1
- First order: s−1
- Second order: dm3 mol−1 s−1
- A larger k means a faster reaction.
Order of Reaction
The order with respect to a reactant is the power to which its concentration is raised in the experimentally determined rate law. The overall order is the sum of these powers.
Order vs molecularity
- Order — experimental, can be 0, fractional, or negative; refers to the overall reaction.
- Molecularity — theoretical, always a positive integer (1, 2, 3); refers to a single elementary step (number of species colliding).
Reading the order off a rate law
| Order in A | Rate law | Doubling [A] does what to the rate? | Units of k |
|---|---|---|---|
| Zero | rate = k | No change | mol dm−3 s−1 |
| First | rate = k[A] | Doubles it (×2) | s−1 |
| Second | rate = k[A]2 | Quadruples it (×4) | dm3 mol−1 s−1 |
This doubling test is how order is found experimentally: change one reactant's concentration, hold the rest fixed, and see what the rate does.
Activation Energy
The activation energy Ea is the minimum energy that colliding molecules must possess (above their average) for a successful reaction. It is the height of the energy barrier between reactants and products on a potential-energy diagram. The peak of that barrier is the transition state (activated complex).
Relating Ea and the activated complex to rate
Only collisions that are both energetic enough (energy ≥ Ea) and correctly oriented form the activated complex and go on to products. So:
- A high Ea means few molecules can surmount the barrier — the reaction is slow.
- A low Ea means a large fraction of collisions succeed — the reaction is fast.
- Raising the temperature does not change Ea; it increases the fraction of molecules that possess it, so the rate (and k) rises.
- The activated complex sits at the peak of the energy barrier. It is unstable and transient, and can fall forward to products or back to reactants.
Catalysts and Ea
A catalyst provides an alternative pathway with a lower activation energy. It speeds up forward and reverse reactions equally, so equilibrium position is unaffected; only the speed at which equilibrium is reached changes. A catalyst is recovered chemically unchanged at the end of the reaction.
Factors Affecting Rate of Reaction
- Concentration of reactants: increasing [reactants] increases collision frequency, raising rate (up to the order). For gases, increasing pressure does the same thing.
- Temperature: a 10 °C rise typically doubles the rate. Higher T → higher fraction of molecules with energy ≥ Ea, and more frequent collisions.
- Catalyst: lowers Ea, providing an alternative pathway. Heterogeneous (Pt, Ni, V2O5) or homogeneous (acids in ester hydrolysis).
- Surface area (heterogeneous reactions): finer particles → greater contact area → faster reaction. Powdered Mg burns faster than ribbon.
- Nature of reactants: ionic reactions in solution are nearly instantaneous; covalent bond rearrangements (organic) are slower.
- Light (photochemical reactions): H2 + Cl2 reacts violently in sunlight but very slowly in the dark.
Worked MCQs
Five MCQs that capture the high-yield testing patterns for this chapter. Read the explanation even when you get the answer right — it's where the deeper concept lives.
Q1. Which of the following is true about a catalyst?
A catalyst lowers Ea by offering an alternative path. It speeds up both forward and reverse reactions equally, so the equilibrium position is unchanged, and it is recovered chemically unchanged at the end.
Q2. The rate constant k of a reaction is changed by:
k is independent of concentration — changing [reactant] changes the rate through the rate law, not k itself. k depends on temperature (and on whether a catalyst is present), which is why the same reaction has a different k at a different temperature.
Q3. For a reaction with rate law rate = k[A]2[B], the overall order is:
Overall order is the sum of the powers in the experimentally determined rate law. Here 2 + 1 = 3, so the reaction is third order overall (second order in A and first order in B).
Q4. Increasing the temperature of a reaction by 10 °C usually:
A 10 °C rise typically doubles the rate because a much larger fraction of molecules now have energy above Ea, and collisions are more frequent. Ea itself is unchanged — only the rate constant k changes.
Q5. The units of the rate constant for a first-order reaction are:
For first order, rate = k[A]. Rate has units mol dm−3 s−1 and [A] has mol dm−3, so k has units of s−1. (Zero order: mol dm−3 s−1; second order: dm3 mol−1 s−1.)
Quick Recap
- Rate = −d[reactant]/dt = +d[product]/dt; units mol dm−3 s−1.
- Rate law: rate = k[A]m[B]n; m + n = overall order (experimental).
- Order vs molecularity: order is experimental and can be 0/fractional; molecularity is theoretical, integral, for elementary steps only.
- k depends on temperature and catalyst only — never on concentration; its units follow the overall order.
- Ea = minimum energy for a successful collision; the activated complex is the unstable species at the top of the barrier. Lower Ea → faster reaction; +10 °C ~ doubles the rate.
- Catalyst ↓ Ea via alternative pathway; equilibrium unchanged.
- Factors: concentration, temperature, catalyst, surface area, nature of reactants, light.